Chapter 3B
Modern Atomic
Theory
1
CHAPTER OUTLINE

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
Waves
Electromagnetic Radiation
Dual Nature of Light
Bohr Model of Atom
Quantum Mechanical Model of Atom
Electron Configuration
Electron Configuration & Periodic Table
Abbreviated Electron Configuration
2
WAVES
 All
Wavelength
waves are(λ)
characterized
is the distance
bybetween
wavelength,
any 2
wavelength
frequency and
successive
crests
speed.
or troughs.
(measuredwavelength
from
peak to(measured
peak)
from
trough to trough)
3
10.1
WAVES
 Frequency (nu,) is the number of waves
produced per unit time.
 Wavelength and frequency are inversely
proportional.
As wavelength of a wave increases
its frequency decreases
inversely
 proportional
Speed tells how fast waves travel through space.
4
10.1
ELECTROMAGNETIC
RADIATION
 Energy travels through space as
electromagnetic radiation. This radiation
takes many forms, such as sunlight,
microwaves, radio waves, etc.
 In vacuum, all electromagnetic waves travel at
the speed of light (3.00 x 108 m/s), and differ
from each other in their frequency and
wavelength.
5
ELECTROMAGNETIC
RADIATION
classification
electromagnetic
 The
These
waves rangeoffrom
-rays (shortwaves
λ, high f)
Long
according
to
radio
waves
to their
(long
frequency
λ, low
f).is called
Short
wavelength
wavelength
electromagnetic
spectrum.
High frequency
Low frequency
6
ELECTROMAGNETIC
RADIATION
Visible
light is waves
a
Infrared
have
X-rays have longer
small part of the EM
longer λ but lower 
λ but lower  than spectrum
than visible light
-rays
7
10.2
DUAL NATURE
OF LIGHT
 When white light is passed through a glass
prism, it is dispersed into a spectrum of colors.
 This is evidence of the wave nature of light.
8
DUAL NATURE
OF LIGHT

have
much evidence
that
light
RedScientists
light has also
longer
wavelength
and less
energy
act as a stream of tiny particles, called
thanbeams
blue light
photons.
A photon of
red light
A photon of
blue light
9
DUAL NATURE
OF LIGHT
 Scientists, therefore, use both the wave and
particle models for explaining light. This is
referred to as the wave-particle nature of light.
 Flatland Video
 Scientists also discovered that when atoms are
energized at high temperatures or by high
voltage, they can radiate light. Neon lights are
an example of this property of atoms.
10
ATOMIC LINE
SPECTRUM
 These
Each element
When
lines
the light
indicate
possesses
fromthat
theaatom
light
unique
isis formed
placed
line spectrum
only at
through
certain
that
canwavelengths
abeprism,
used to
a series
identify
and frequencies
of it.
brightly colored
that
lights, calledtoa specific
line spectrum
correspond
colors.is formed.
Each line represents a particular  and 
11
BOHR MODEL
OF ATOM
 Bohr’s
In thisBohr,
Neils
model,
model
a Danish
of
thethe physicist, studied the
hydrogen
atom
electrons
consisted
could
atom of
only
extensively, and developed a
model for
electrons
occupy
particular
orbiting
the atomthe
energy
that was able to explain
the lineand
nucleus
levels,
spectrum.
at different
could “jump”
distances
to
higher levels
from the
by
nucleus, called
energy
absorbing
energy.
levels.
12
BOHR MODEL
OF ATOM
 The lowest energy level is
called ground state, and
the higher energy levels
are called excited states.
 When electrons absorb
energy through heating
or electricity, they move
to higher energy levels
and become excited.
energy
13
BOHR MODEL
OF ATOM
 When excited electrons
return to Lower
the ground
state,
energy
energy is emitted
as a
transition
photon of light
is released.
give off
red
 The color (wavelength)
of
light
the lightHigher
emittedenergy
is
determined
by thegive
transition
difference
energy
offinblue
light
between the two states
(excited and ground).
14
BOHR MODEL
OF ATOM
 The line spectrum is produced by many of these
transitions between excited and ground states.
 Bohr’s model was able to successfully explain
the hydrogen atom, but could not be applied to
larger atoms.
 Quantum Mechanics & Structure of Atom
15
QUANTUM MECHANICAL
MODEL OF ATOM
 In 1926 Erwin Shrödinger created a mathematical
model that showed electrons as both particles and
waves. This model was called the quantum
mechanical model.
 Double-Slit Experiment
 This model predicted electrons to be located in a
probability region called orbitals.
 An orbital is defined as a region around the
nucleus where there is a high probability of
finding an electron.
16
QUANTUM MECHANICAL
MODEL OF ATOM
 Based on this model, there
are discrete principal
energy levels within the
atom.
 Principal
levelsthe
are
As energy
n increases,
designatedenergy
by n. of the
 The electrons
inincreases
an atom
electron
can exist in any principal
energy level.
17
QUANTUM MECHANICAL
MODEL OF ATOM
 Each principal energy level
is subdivided into sublevels.
 The sublevels are
designated by the letters
s, p, d and f.
 As n increases, the number
of sublevels increases.
18
10.7, 10.8
QUANTUM MECHANICAL
MODEL OF ATOM
 Within
The number
the sublevels,
of orbitals
thewithin
electrons
the sublevels
are located in
orbitals.
vary
withThe
theirorbitals
type. are also designated by the
letters s, p, d and f.
s sublevel = 1 orbital = 2 electrons
p sublevel = 3 orbitals = 6 electrons
d sublevel = 5 orbitals = 10 electrons
f sublevel = 7 orbitals = 14 electrons
An orbital can hold a maximum of 2 electrons
19
ELECTRON
CONFIGURATION
 Similarities of behavior in the periodic table are
due to the similarities in the electron
arrangement of the atoms. This arrangement is
called electron configuration.
 The modern model of the atom describes the
electron cloud consisting of separate energy
levels, each containing a fixed number of
electrons.
 Each orbital can be occupied by no more than 2
electrons, each with opposite spins.
20
ELECTRON
CONFIGURATION
 The electrons occupy the orbitals from the lowest
energy level to the highest level.
 The energy of the orbitals on any level are in the
following order: s < p < d < f.
 Each orbital on a sublevel must be occupied by a
single electron before a second electron enters. For
example, all three p orbitals must contain one
electron before a second electron enters a p orbital
(Hund’s Rule).
 Visualizing Orbitals
21
ELECTRON
CONFIGURATION
 Electron configurations can be written as:
2
Principal
energy level
6
p
Number of
electrons in
orbitals
Type of
orbital
22
ELECTRON
CONFIGURATION
 Another notation, called the orbital notation is
shown below:
Electrons in
orbital with
opposing spins
Principal
energy level
Type of
orbital
1s
23
ELECTRON
CONFIGURATION
H
↑
1s1
1s
Hydrogen has 1 electron. It will occupy the orbital of
lowest energy which is the 1s.
He
↑↓
1s2
1s
Helium has two electrons. Both helium electrons occupy the
1s orbital with opposite spins.
24
ELECTRON
CONFIGURATION
Li
↑↓
1s
↑
1s22s1
2s
The 1s orbital is filled. Lithium’s third electron will enter the
2s orbital.
Be
↑↓
↑↓
1s22s2
1s
2s
The 2s orbital fills upon the addition of beryllium’s third and
fourth electrons.
25
ELECTRON
CONFIGURATION
B
↑↓
↑↓
1s
2s
↑
1s22s22p1
2p
Boron has the first p electron. The three 2p orbitals have the
same energy. It does not matter which orbital fills first.
C
↑↓
↑↓
1s
2s
↑↓ ↑
1s22s22p2
2p
The second p electron of carbon enters a different p orbital
than the first p due to Hund’s Rule.
26
ELECTRON
CONFIGURATION
N ↑↓
↑↓
1s
2s
↑
↑
↑
1s22s22p3
2p
The third p electron of nitrogen enters a different p orbital than
its first two p electrons due to Hund’s Rule.
O
↑↓
↑↓
1s
2s
↑↓ ↑
↑
1s22s22p4
2p
The last p electron of oxygen pairs opposite of another since
each orbital has an electron in it and Hund’s Rule is satisfied.
27
ELECTRON
CONFIGURATION
F
↑↓
↑↓
1s
2s
↑↓ ↑↓ ↑
1s22s22p5
2p
Two of the p electrons for fluorine pair up with other electrons
in the p orbitals.
Ne
↑↓
↑↓
↑↓ ↑↓ ↑↓
1s
2s
2p
1s22s22p6
The last p electron for neon pairs up with the last lone electron
and completely fills the 2nd energy level.
28
ELECTRON
CONFIGURATION
 As electrons occupy the 3rd energy level and
higher, some anomalies occur in the order of the
energy of the orbitals.
 Knowledge of these anomalies is important in
order to determine the correct electron
configuration for the atoms.
 The following study aid is used by beginning
students to remember these exceptions to the
order of orbital energies.
29
ELECTRON
CONFIGURATION
1s 2s 2p 3s 3p 4s 3d 4p 5s 4d 5p 6s
30
ELECTRON CONFIG.
& PERIODIC TABLE
 The horizontal rows in the periodic table are
called periods. The period number corresponds
to the number of energy levels that are occupied
in that atom.
 The vertical columns in the periodic table are
called groups or families. For the main-group
elements, the group number corresponds to the
number of electrons in the outermost filled
energy level (valence electrons).
31
ELECTRON CONFIG.
& PERIODIC TABLE
energy
4One
energy
3 energy
level levels
levels
32
ELECTRON CONFIG.
& PERIODIC TABLE
3 valence
1 valence
5 valence
electrons
electron
electrons
33
ELECTRON CONFIG.
& PERIODIC TABLE
 The
Notevalence
that elements
electrons
in the
configuration
same groupfor
have
the
elementselectron
similar
in periods
configurations.
1-3 are shown below.
34
10.15
ELECTRON CONFIG.
& PERIODIC TABLE
Arrangement of orbitals in the periodic table
35
10.16
ELECTRON CONFIG.
& PERIODIC TABLE
d orbital numbers are 1 less
than the period number
36
10.16
ELECTRON CONFIG.
& PERIODIC TABLE
f orbital numbers are 2 less
than the period number
37
10.16
ELECTRON CONFIG.
& PERIODIC TABLE
 The electrons in an atom fill from the lowest to
the highest orbitals.
 The knowledge of the location of the orbitals on
the periodic table can greatly help the writing of
electron configurations for large atoms.
 The energy order of the sublevels is shown next.
Note that some anomalies occur in the energy
level of “d” and “f” sublevels.
38
10.15
ELECTRON CONFIG.
& PERIODIC TABLE
39
10.15
Example 1:
Use the periodic table to write complete electron
configuration for phosphorus.
P
Z = 15
10
electrons
5 electrons
used
remaining
1s2 2s2 2p6 3s2 3p3
Core
electrons
Valence
electrons
40
Example 2:
Draw an orbital notation diagram for the last incomplete level of chlorine and determine the number
of unpaired electrons.
3p
3s
41
Example 2:
Cl
↑↓
↑↓ ↑↓ ↑
3s
3p
One
unpaired
electron
42
ABBREVIATED
ELECTRON CONFIG.
 When writing electron configurations for larger
atoms, an abbreviated configuration is used.
 In writing this configuration, the non-valence
(core) electrons are summarized by writing the
symbol of the noble gas prior to the element in
brackets followed by configuration of the
valence electrons.
43
ABBREVIATED
ELECTRON CONFIG.
K
Z = 19
1s22s22p63s23p6 4s1
core[Ar]
electrons
4s1
Previous
noble gas
valence
electron
44
ABBREVIATED
ELECTRON CONFIG.
Br
Z = 35
1s22s22p63s23p6 4s2 3d10 4p5
core[Ar]
electrons
4s23d104p5
valence
electrons
45
Example 3:
Write abbreviated electron configurations for
each element listed below:
Fe
Z = 26
8 electrons
18
electrons
620electrons
electrons
remaining
used
remaining
used
[Ar] 4s2 3d6
46
Example 3:
Write abbreviated electron configurations for
each element listed below:
Sb
Z = 51
15
36electrons
electrons
38 electrons
13
48
3 electrons
electrons
remaining
used
remaining
used
remaining
used
[Kr] 5s2 4d10 5p3
5
valence
electrons
47
TRENDS IN
PERIODIC PROPERTIES
 The electron configuration of atoms are an
important factor in the physical and chemical
properties of the elements.
 Some of these properties include: atomic size,
ionization energy and metallic character.
 These properties are commonly known as
periodic properties and increase or decrease
across a period or group, and are repeated in
each successive period or group.
48
ATOMIC SIZE
 The size of the atom is determined by its atomic
radius, which is the distance of the valence
electron from the nucleus.
 For each group of the representative elements,
the atomic size increases going down the group,
because the valence electrons from each energy
level are further from the nucleus.
49
ATOMIC SIZE
50
ATOMIC SIZE
 The atomic radius of the representative elements
are affected by the number of protons in the
nucleus (nuclear charge).
 For elements going across a period, the atomic
size decreases because the increased nuclear
charge of each atom pulls the electrons closer to
the nucleus, making it smaller.
51
ATOMIC SIZE
52
IONIZATION
ENERGY
 The ionization energy is the energy required to
remove a valence electron from the atom in a
gaseous state.
 When an electron is removed from an atom, a
cation (+ ion) with a 1+ charge is formed.
Na (g) + IE
Na+ + e-
53
IONIZATION
ENERGY
 The ionization energy
decreases going down a
group, because less energy
is required toLarger
removeatom
an
Less
IE
electron from the
outer
shell since it is further
from the nucleus.
54
IONIZATION
ENERGY
 Going across a period,
the ionization energy
increases because the
increased nuclear
charge of the atom
holds the valence
electrons more tightly
and therefore it is more
difficult to remove.
55
IONIZATION
ENERGY
 In general, the ionization
energy is low for metals
and high for non-metals.
 Review of ionization
energies of elements in
periods 2-4 indicate
some anomalies to the
general increasing trend.
56
IONIZATION
ENERGY
 These anomalies are caused by more stable
More
stable
electron configurations
ofstable
the atoms in groups 2
More
(1/2and
filled)
(complete “s” sublevel)
group 5 (half-filled “p”
Higher
IE
Higher
IE in their ionization
sublevels) that cause an
increase
energy compared to the next element.
Be
1s2 2s2
N
1s2 2s2 2p3
B
1s2 2s2 2p1
O
1s2 2s2 2p4
57
METALLIC
CHARACTER
 Metallic character is the ability of an atom to lose
electrons easily.
 This character is more prevalent in the elements
on the left side of the periodic table (metals), and
decreases going across a period and increases for
elements going down a group.
58
METALLIC
CHARACTER
Most
metallic
elements
Least
metallic
elements
59
Example 1:
Select the element in each pair with the larger
atomic radius:
Li
or
K
Larger due
to more
energy levels
60
Example 1:
Select the element in each pair with the larger
atomic radius:
K
or
Br
Larger due to
less nuclear
charge
61
Example 1:
Select the element in each pair with the larger
atomic radius:
P
or
Cl
Larger due to
less nuclear
charge
62
Example 2:
Indicate the element in each set that has the
higher ionization energy and explain your choice:
K
or
Na
Higher IE
due to less
energy levels
63
Example 2:
Indicate the element in each set that has the
higher ionization energy and explain your choice:
Mg
or
Cl
Higher IE due
to more
nuclear charge
64
Example 2:
Indicate the element in each set that has the higher
ionization energy and explain your choice:
F
N
or
C
Highest IE due
to most
nuclear charge
65
THE END
66